Chapter 6: Representing Molecules – Lewis Structures, Bonding, and Polarity

Representing Molecules

The Octet Rule and Lewis Structures

Understanding how atoms bond and share electrons is fundamental to representing molecules in chemistry. The octet rule and Lewis structures provide a framework for visualizing the arrangement of electrons in molecules.

• Covalent Bond: A chemical bond where two or more electrons are shared by two atoms. This sharing allows atoms to achieve a stable electron configuration.

• Bonding Pairs: Electrons shared between atoms, forming a bond.

• Lone Pairs: Non-shared electrons that remain on an atom.

Example: In the F2 molecule, each fluorine atom shares one electron to form a single covalent bond, achieving an octet.

Octet Rule: Atoms (except hydrogen) form bonds until they have 8 valence electrons, resembling the electron configuration of noble gases. Hydrogen follows the duet rule, achieving 2 electrons.

• Lewis Structure of H2: Each hydrogen atom shares one electron, resulting in a single covalent bond and a duet configuration.

• Lewis Structure of H2O: Oxygen shares electrons with two hydrogens, forming two single covalent bonds and two lone pairs, achieving an octet for oxygen.

Multiple Bonds

Atoms can share more than one pair of electrons, forming double or triple bonds. This allows for the sharing of extra electrons when needed for the central atom to achieve an octet.

• Double Bond: Two atoms share two pairs of electrons. Example: O2 and CO2.

• Triple Bond: Two atoms share three pairs of electrons. Example: N2 and HCN.

Example: In CO2, each oxygen forms a double bond with carbon, allowing all atoms to achieve an octet.

Bond Length

Bond length is the average distance between the nuclei of two covalently bonded atoms. The type of bond affects the bond length:

• Triple bond < Double bond < Single bond (for the same atoms)

Example: H2 has a bond length of 74 pm, while HI has a bond length of 161 pm.

Bond Type	Bond Length (pm)
C—H	109
C—C	154
C=C	134
C≡C	120
C—N	147
C=N	129
C≡N	116
N—O	140
N=O	120
O—H	96

Electronegativity and Polarity

Polar Covalent Bonds

When atoms in a bond have different electronegativities, electrons are shared unequally, resulting in a polar covalent bond. The more electronegative atom carries a partial negative charge (δ-), while the less electronegative atom carries a partial positive charge (δ+).

• Example: In HF, fluorine is more electronegative and attracts electrons more strongly, creating a dipole.

Electronegativity

Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond. The difference in electronegativity between atoms determines the bond's polarity.

• Fluorine (F) is the most electronegative element.

• Electronegativity increases across a period and decreases down a group in the periodic table.

Example: In H—F, fluorine attracts electrons more strongly than hydrogen, resulting in a polar bond.

Classification of Bonds Based on Electronegativity Difference

Bonds are classified by the difference in electronegativity between the atoms:

• Non-polar covalent: Equal sharing of electrons (e.g., H2, Cl2).

• Polar covalent: Unequal sharing, partial transfer of electrons (e.g., H—F, C—F).

• Ionic: Full transfer of electrons (e.g., NaCl, KF).

Percent Ionic Character measures the polarity of a bond. 100% ionic means complete electron transfer; 100% covalent means equal sharing.

Bond	Type	Polarity Ranking
H—H	Non-polar covalent	1 (least polar)
NaCl	Ionic	4 (most polar)
C—H	Polar covalent	2
KF	Ionic	4 (most polar)
Cl—Cl	Non-polar covalent	1 (least polar)
C—F	Polar covalent	3

Example: NaCl and KF are ionic and most polar; H—H and Cl—Cl are non-polar covalent and least polar.

Additional info:

• Lewis structures are essential for predicting molecular geometry, reactivity, and physical properties.

• Bond length and bond polarity influence molecular interactions and chemical behavior.